Materials Included In Kit

Additional Materials Required

Safety Precautions

Hydrochloric acid solution is toxic and corrosive to eyes and skin tissue. Sodium hydroxide solution is a hazardous and corrosive liquid; it is dangerous to eyes and can cause skin burns. Ferric nitrate solution is a possible skin and body tissue irritant; it will also stain clothes and skin. Potassium thiocyanate is toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Clean up all chemical spills immediately. Wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling, and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. All final solutions can be disposed of according to Flinn Suggested Disposal Method #26b.

Teacher Tips

• Enough materials are provided in this kit for 30 students working in pairs, or for 15 groups of students. The experimental work can reasonably be completed in one 50-minute class period. The prelab activity may be assigned separately as preparation for lab, or it can be used as the basis of a cooperative class discussion.
• Enough disposable pipets have been provided to give each pair of students 10 pipets. Encourage your students to label their pipets and to rinse them between parts of the experiment to avoid contamination and waste.
• Several Flinn demonstration kits have been developed to reinforce predictions made using Le Chatelier’s Principle (Pink and Blue: A Colorful Chemical Balancing Act, AP8471, and Hot and Cold Equilibrium: A Demonstration of Le Chatelier’s Principle, AP5938). These demonstration kits utilize the complex-ion equilibrium involving CoCl42– and Co(H2O)62+. This is also the equilibrium observed in moisture-sensitive Hydrion Humidicator paper (AP4656).
• The effect of CO2 in exhaled air on the color of bromthymol blue indicator provides an interesting extension of acid–base equilibria involving indicators. Fill three clear plastic food cups (not lab glassware!) with tap water. Add a few drops of lemon juice or vinegar to the first cup. Have a student blow through a fresh, clean straw into the water in the second cup for about 3 minutes. Immediately afterwards, add 5 drops of bromthymol blue indicator solution to each cup and compare the colors of the solutions. The water that has been blown into becomes acidic due to the reaction of CO2 with water to give carbonic acid (H2CO3).
• The green transition color (step 6) in the equilibrium reaction of bromcresol green indicator solution is easy to overshoot. Students should add HCl carefully drop by drop, taking care to gently shake the solution between drops. Toward the end, when the color begins to change more slowly between drops, have them try to add half a drop at a time. (Squeeze out a small amount from the pipet, press the pipet tip against the side of the test tube to dislodge the half-drop, and then swirl the test-tube contents to rinse the half-drop into solution.)
• The effect of acid on the solubility equilibrium of magnesium hydroxide can be used to illustrate the action and effectiveness of Milk of Magnesia, a popular antacid. Milk of Magnesia is a suspension of solid magnesium hydroxide in water. The suspension dissolves as required in the stomach to combat excess acidity.
• Challenge your students to apply their understanding of LeChâtelier’s Principle to explain the effects of high altitudes on humans. The protein hemoglobin (Hb), which is responsible for the transport of oxygen in the blood, contains four iron ions that are able to bind oxygen molecules. This must be a reversible reaction, since the hemoglobin must be able to release the oxygen molecules in cells and body tissue. {13366_Tips_Reaction_1} At high altitudes, where the concentration of oxygen in the air is lower, this equilibrium is shifted in the reverse direction, and less oxygen is available in the bloodstream to be transported to the cells. The physical symptoms of the reduced oxygen availability are fatigue and dizziness. The human body, however, is marvelous in its adaptability. People who live or train at high altitudes compensate for the reduced oxygen supply by synthesizing more hemoglobin molecules. Use Le Chatelier’s Principle to explain the effect of increasing the hemoglobin concentration on the position of equilibrium for the above reaction.

Answers to Prelab Questions

1. The following statements express common misunderstandings concerning the nature of equilibrium. Modify each statement so that it correctly describes the true nature of chemical equilibrium. a. At equilibrium no more reactants are transformed into products. b. Reactions at equilibrium have equal concentrations of reactants and products. Answer: a. At equilibrium, the rate at which reactants are transformed into products equals the rate at which products are transformed into reactants. b. Reactions at equilibrium have constant concentrations of reactants and products. 2. Acetic acid (CH3COOH) is the principal ingredient in vinegar. Consider the following reversible reaction of acetic acid at equilibrium: {13366_PreLabAnswers_Reaction_1} a. Use Le Chatelier’s Principle to predict the effect of adding excess sodium acetate (Na+CH3COO–) on the concentration of H+ present at equilibrium. b. Why will the addition of OH– ions (in the form of sodium hydroxide) cause the concentration of CH3COO– to increase? Answer: a. Adding excess product, in the form of acetate ion, shifts the equilibrium in the reverse direction to make more reactants and thus reduce the excess product concentration. This effect reduces the concentration of H+ present at equilibrium. b. OH– ions added to the solution react with and neutralize some of the H+ product ions present at equilibrium. According to Le Chatelier’s Principle, removing one of the products of the reaction shifts the equilibrium in the forward direction to make more products. This restoring effect of equilibrium causes the concentration of acetate ion to increase. 3. Papers coated with cobalt chloride powder are sold commercially as moisture-sensitive test strips to estimate relative humidity levels between 20 and 80 percent in air. The following equilibrium reaction takes place with water: {13366_PreLabAnswers_Reaction_2} a. What color will the paper be when the relative humidity is low (20%) versus high (80%)? b. A color chart is available to estimate intermediate humidity levels. Predict the transition colors that might be observed when the relative humidity levels are 40 and 60%, respectively. Answer: a. The paper is blue when the relative humidity is low (20%), because there is not enough water present to react with the CoCl2. When the relative humidity is high (80%), and the amount of water needed to react with CoCl2 is more readily available, the paper turns pink. b. At intermediate humidity levels both reactants and products are present at equilibrium, and intermediate or transition colors are observed midway between blue and pink. Thus, the color changes gradually from blue to purple to lavender to pink.

Sample Data

Part A. Acid–Base Equilibrium of Bromcresol Green {13366_Data_Table_1} Part B. Solubility Equilibrium of Magnesium Hydroxide {13366_Data_Table_2} Part C. Complex-Ion Equilibrium of Iron(III) and Thiocyanate Ion {13366_Data_Table_3}

Answers to Questions

1. Write the chemical equation for the reversible chemical reaction that occurs when the indicator bromcresol green (HIn) is dissolved in water (Part A). Answer: {13366_Answers_Reaction_1} 2. Use Le Chatelier’s Principle to explain the color changes observed upon addition of HCl and NaOH, respectively, to a solution of bromcresol green in water (steps 3 and 4). Answer: Adding HCl increases the concentration of product and thus shifts the equilibrium in the reverse direction to use up the excess product. This increases the amount of the acid form, HIn, which is yellow, and so a color change from blue to yellow is observed. Adding NaOH has the opposite effect. It reacts with and neutralizes H+ ions in solution, thus decreasing the concentration of one of the products. Removing a product shifts the equilibrium in the forward direction and thus increases the concentration of the basic form, In–, which is blue. 3. What is the likely composition of the indicator solution when the intermediate or transition color is observed in step 6? How does this observation provide visual proof of the idea that not all reactions “go to completion”? Explain. Answer: The intermediate (green) color of the indicator solution suggests that at this point approximately half of the total available indicator in solution is present in the form of uncharged HIn molecules (yellow) and half in the form of In– ions (blue). The green color, midway between blue and yellow, is visual proof that at this point both reactants and products must be present at equilibrium, that is, the reaction does not go to completion. 4. Write the chemical equation for the reversible chemical reaction that occurs when magnesium hydroxide is dissolved in water (Part B). Answer: {13366_Answers_Reaction_2} 5. Use Le Chatelier’s Principle to explain the changes observed upon addition of HCl and NaOH, respectively, to a saturated solution of magnesium hydroxide in water (steps 11 and 12). Answer: Adding HCl to a saturated solution (solid in equilibrium with dissolved ions) of Mg(OH)2 causes more solid to dissolve. HCl shifts the equilibrium to the right due to the reaction of H+ ions with OH– ions (one of the products of the reaction). Removing a product from a system at equilibrium shifts a reaction in the forward direction, to make more product! Adding NaOH had the opposite effect; more solid precipitated out of solution. Adding product to a system at equilibrium shifts a reaction in the reverse direction, to use up the excess product. 6. Predict the direction in which the solubility equilibrium should shift when EDTA is added to a saturated solution of magnesium hydroxide in water (step 13). Do your observations support this prediction? Answer: EDTA reacts quantitatively with Mg2+ ions, effectively removing them as products from the solubility equilibrium of magnesium hydroxide. Removing a product should shift the reaction in the forward direction, toward dissolving more solid magnesium hydroxide. This is consistent with the observations that the solid dissolved as EDTA was added to the solution. 7. Write the chemical equation for the reversible chemical reaction that occurs when iron(III) and thiocyanate ion are mixed together in aqueous solution. Answers: {13366_Answers_Reaction_2} 8. How was the composition of the solution (relative amounts of reactants and products) affected when the solution was cooled (step 19)? When the solution was heated (step 20)? Answer: Cooling the solution resulted in a deeper orange-red color due to the formation of more complex-ion product, which is red. Heating the solution resulted in the opposite effect—the amount of the red complex-ion product decreased substantially and the solution reverted to a pale yellow color. 9. Based on the observed effect of temperature on the position of equilibrium, is the forward reaction for this reaction endothermic or exothermic? Explain, using Le Chatelier’s Principle. Answer: The forward reaction, formation of the FeSCN2+ complex-ion product, must be exothermic. According to Le Chatelier’s Principle, increasing the temperature of an exothermic chemical reaction shifts the reaction in the reverse direction, that is, in the direction in which the excess heat can be consumed. 10. Use Le Chatelier’s Principle to explain the color changes observed upon addition of excess Fe3+ and SCN–, respectively, to an equilibrium mixture of reactants and products in this complex-ion reaction (steps 21 and 22). Answer: Addition of excess reactants, both Fe3+ and SCN–, resulted in the solutions turning a deep, blood red color. According to Le Chatelier’s Principle, adding excess reactant to a system at equilibrium shifts the reaction in the forward direction, to make more product (red) and thus use up the excess reactants. 11. After observing the effect of NaH2PO4 on the equilibrium mixture (step 23), a student was skeptical that both Fe3+ and SCN– were still present in solution. Can you think of additional experimental steps that could be done to prove that both reactants are still present at this point? Answer: The solution was colorless after NaH2PO4 was added. Try adding more Fe3+ to the solution again at this point. If SCN– is still present, it will react with the added Fe3+, and, indeed, that is what is observed—the solution turns deep red again. The same holds true if more SCN– is added to the colorless solution.

References

Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry; University of Wisconsin Press: Madison, 1983; pp 338–343. Source Book Version 1.0, Orna, M. V., Schreck, J. O., and Heikkinen, H., Eds.; ChemSource: New Rochelle, NY, 1994; Vol. 2, pp 2–28. Vonderbrink, Sally Ann Laboratory Experiments for Advanced Placement Chemistry; Flinn Scientific: Batavia, IL, 1995; pp 109–113.

Introduction

The word equilibrium has two roots: æqui, meaning equal, and libra, meaning weight or balance. Our physical sense of equilibrium suggests a condition of equal balance of opposing forces. How does this physical sense of equilibrium translate to chemical equilibrium? The purpose of this lab is to explore the nature and consequences of equilibrium in a variety of chemical reactions.

Concepts

• Reversible chemical reaction
• Chemical equilibrium
• Le Chatelier’s Principle
• Endothermic vs. exothermic reactions

Background

In our first exposure to the nature of chemical change, we learn to distinguish between chemical and physical changes by emphasizing that physical changes can be easily reversed. In contrast, early examples of chemical change often feature reactions that cannot be reversed, such as burning wood, rusting metal, or spoiling food. A closer look at chemical change, however, reveals that many chemical reactions are indeed reversible. Consider the following example of a reversible chemical reaction. Nitrogen and hydrogen gas react under high-pressure conditions to make ammonia. At relatively high temperatures, however, ammonia decomposes to reform the nitrogen and hydrogen starting materials. The reaction can go both ways! This reversible chemical reaction is represented symbolically using double arrows, as shown in Equation 1. {13366_Background_Equation_1} What happens when nitrogen gas and hydrogen gas are allowed to react? As ammonia is produced, the amounts of nitrogen and hydrogen decrease and the amount of ammonia begins to increase. As the concentration of ammonia increases, however, the reverse reaction begins to take place at a significant rate as well. Eventually, the amount of ammonia consumed due to the reverse reaction becomes equal to the amount of ammonia formed in the forward reaction. At this point, no further changes are observed in the amounts of either the reactants or product. This is the position of chemical equilibrium. Chemical equilibrium is defined as the state where the concentrations of reactants and products remain constant with time. This does not mean that both reactions have stopped. Both the forward and reverse reactions are still taking place, but at equal rates, so the amounts of reactants and products remain constant. Chemical equilibrium represents an equal balance of opposing forces due to the forward and reverse reactions. Le Chatelier’s Principle What happens when the balance is disturbed—due to changes in reactant and product concentrations, temperature, or pressure, for example? Le Chatelier’s Principle provides an intuitive explanation for how balance can be restored: “If an equilibrium system is subjected to a stress, the system will react in such a way as to remove the stress.” To remove a stress, one of two things can happen. The reaction can shift in the forward direction to make more products, thus using up reactants. Alternatively, the reaction can shift in the reverse direction to re-form the reactants, thus using up products. Adding a reactant to an equilibrium mixture of reactants and products always shifts a reversible chemical reaction in the forward direction, to use up the excess reactant and make more products. Removing a product from a mixture already at equilibrium has the same effect. Conversely, adding a product or removing a reactant has the opposite effect, shifting a reaction in the reverse direction to use up the excess product and make more reactants. The effect of temperature on a system at equilibrium depends on whether a reaction is endothermic or exothermic. If a reaction is endothermic, heat may be considered as a reactant in the forward reaction. (Endothermic reactions require heat in order to proceed.) Increasing the temperature of an endothermic, reversible chemical reaction shifts the position of equilibrium in the forward direction, to use up the excess heat and make more products. In the case of an exothermic chemical reaction, heat may be thought of as a product of the forward reaction. Increasing the temperature of an exothermic, reversible chemical reaction shifts the position of equilibrium in the reverse direction, in order to remove the excess heat represented by the higher temperature. The ammonia synthesis reaction is exothermic (Equation 2). Increasing the temperature of this reaction shifts the equilibrium in the reverse direction, to use up some of the excess heat. Ammonia decomposition becomes more significant at higher temperatures. {13366_Background_Equation_2} Equilibria governing reactions in aqueous solutions play a vital role in environmental and biological chemistry. The pH and alkalinity of drinking water depend on acid–base equilibria. The factors that control the movement of solutes in groundwater or surface water depend on solubility equilibria. The binding of oxygen to hemoglobin depends on complex-ion equilibria involving iron(II) ions bound to porphyrin complexes in the protein. Overview of the Experiments The properties of an acid–base indicator illustrate the nature of acid–base equilibria. An indicator is an intensely colored compound that can gain or lose a hydrogen ion to form substances that have different colors. Equation 3 illustrates the reaction of the indicator bromcresol green (HIn). HIn represents the uncharged indicator molecule and In– is the indicator ion after the molecule has lost a hydrogen ion. The position of equilibrium for this reaction depends on the H+ concentration and can easily be shifted in one direction or another by adding or removing H+. In examining the behavior of equilibria involving H+ ions, it is important to remember that OH– reacts irreversibly with H+ to give water (Equation 4). Adding excess OH– to a reversible acid–base reaction has the effect, therefore, of removing H+ from the reaction mixture. {13366_Background_Equation_3} {13366_Background_Equation_4} The solubility of magnesium hydroxide in aqueous solution illustrates the nature of solubility equilibria. A saturated aqueous solution of magnesium hydroxide is prepared by the reaction of magnesium chloride with sodium hydroxide in water (Equation 5). The solubility equilibrium observed in a saturated aqueous solution of magnesium hydroxide involves solid salt on one side and dissolved aqueous ions on the other side (Equation 6). What are the predicted effects of adding excess OH– ion, adding H+ (in the form of HCl), and adding EDTA (a reagent that reacts irreversibly with Mg2+ ion, according to Equation 7)? {13366_Background_Equation_5} {13366_Background_Equation_6} {13366_Background_Equation_7} Reaction of ferric nitrate with potassium thiocyanate illustrates the nature of complex-ion equilibria. Aqueous iron(III) ion (Fe3+) is pale yellow. Adding colorless thiocyanate ion (SCN–) results in the formation of a complex ion in which the anion is tightly bound to the metal ion (Equation 8). The resulting iron–thiocyanate complex ion is a deep red color. The effects of adding excess reactants on the position of equilibrium will be observed. The effect of adding sodium phosphate (NaH2PO4), which undergoes a competing reaction with Fe3+, will also be examined. Finally, the effect of temperature will be studied in order to determine whether the reaction is endothermic or exothermic in the forward direction. {13366_Background_Equation_8}

Materials

Prelab Questions

1. The following statements express common misunderstandings concerning the nature of equilibrium. Modify each statement so that it correctly describes the true nature of chemical equilibrium. a. At equilibrium no more reactants are transformed into products. b. Reactions at equilibrium have equal concentrations of reactants and products. 2. Acetic acid (CH3COOH) is the principal ingredient in vinegar. Consider the following reversible reaction of acetic acid at equilibrium: {13366_PreLab_Reaction_1} a. Use LeChâtelier’s Principle to predict the effect of adding excess sodium acetate (Na+CH3COO–) on the concentration of H+ present at equilibrium. b. Why will the addition of OH– ions (in the form of sodium hydroxide) cause the concentration of CH3COO– to increase? 3. Paper coated with cobalt chloride powder is sold commercially as moisture-sensitive test strips to estimate relative humidity levels between 20 and 80 percent in air. The following equilibrium reaction takes place with water: {13366_PreLab_Reaction_2} a. What color will the paper be when the relative humidity is low (20%) versus high (80%)? b. A color chart is available to estimate intermediate humidity levels. Predict the transition colors that might be observed when the relative humidity levels are 40 and 60%, respectively.

Safety Precautions

Hydrochloric acid solution is toxic and corrosive to eyes and skin tissue. Sodium hydroxide solution is a hazardous and corrosive liquid; it is dangerous to eyes and can cause skin burns. Ferric nitrate solution is a possible skin and body tissue irritant; it will also stain clothes and skin. Potassium thiocyanate is toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Clean up all chemical spills immediately. Wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Preparation 1. Fill two beakers (250- or 400-mL) half-full with water. Prepare an ice-water bath for use in Part C (step 19) and a hot water bath for use in Part C (step 20). Part A. Acid–Base Equilibrium of Bromcresol Green 2. Obtain 2 mL of distilled water in a clean test tube and add 5 drops of 0.04% bromcresol green. Record the color of the solution on the Data Sheet. 3. Add 3 drops of 0.1 M HCl solution. Shake gently. Record the new color of the indicator solution. 4. Add 0.5 M NaOH dropwise to the solution until the original color is restored. Shake gently and record the number of drops of NaOH added and the color of the solution on the Data Sheet. 5. Continue adding 0.5 M NaOH dropwise until a total of 5 drops of NaOH have been added in steps 4 and 5 combined. Can the process be reversed to obtain the color that is intermediate between the colors observed in steps 3 and 4? 6. Add 0.1 M HCl again dropwise very slowly until the solution reaches the intermediate or transition color midway between the two colors. Shake gently between drops to ensure thorough mixing and to avoid overshooting the transition color. Record the number of drops of HCl required and the transition color on the Data Sheet. 7. The contents of the test tube should be disposed of according to teacher instructions. Wash the the test tube and rinse it thoroughly with distilled water after disposal. Part B. Solubility Equilibrium of Magnesium Hydroxide 8. Label two clean test tubes A and B. 9. To each test tube add 1 mL of 1 M MgCl2 solution and 1 drop of universal indicator solution. 10. Add 10 drops of 0.5 M NaOH to each test tube. Record the color and appearance of both solutions on the Data Sheet. 11. Add 1 drop of 0.1 M HCl to the mixture in test tube A. Shake gently and continue adding HCl dropwise until no further changes are observed in the color and appearance of the solution. Record the number of drops of HCl added and the observations on the Data Sheet. 12. Add 10 drops of 0.5 M NaOH to the solution in test tube A. Record the observations on the Data Sheet. 13. To the solution in test tube B add 10 drops of 0.1 M EDTA solution. Record the color and appearance of the solution on the Data Sheet. 14. The contents of the test tubes should be disposed of according to teacher instructions. Wash the test tubes and rinse them thoroughly with distilled water after disposal. Part C. Complex-Ion Equilibrium of Iron(III) and Thiocyanate Ion 15. Prepare a reference solution of FeSCN2+: To a clean 50-mL beaker, add 40 mL of distilled water followed by 1 mL of 0.1-M Fe(NO3)3 solution and 1 mL of 0. 1 M KSCN solution. Mix thoroughly with a glass stirring rod. 16. Label six clean test tubes 1–6. 17. Add 1 mL of the FeSCN2+ reference solution to each test tube 1–6. 18. Add 10 drops of distilled water to test tube 1 and record the color of the solution on the Data Sheet. 19. Add 10 drops of distilled water to test tube 2 and place the sample in an ice-water bath. After 5 minutes, remove the test tube from the ice bath and compare the color of the solution to that of the reference solution in test tube 1. Record the color comparison on the Data Sheet. 20. Add 10 drops of distilled water to test tube 3 and place the sample in a hot water bath at 70–80 °C. After 2–3 minutes remove the tube and record the color of the solution on the Data Sheet. 21. To test tube 4, add 10 drops of 0.1 M Fe(NO3)3. Compare the color of the resulting solution to the solution in test tube 1 and record the color comparison on the Data Sheet. 22. To test tube 5, add 10 drops of 0.1 M KSCN. Compare the color of the resulting solution to the solution in test tube 1 and record the color comparison on the Data Sheet. 23. To test tube 6, add 5 drops of 0.1 M NaH2PO4. Record the color and appearance of the solution on the Data Sheet. 24. Consult your instructor for appropriate disposal procedures.